Notes on teaching Chemistry 126 using Introduction to Molecular Thermodynamics
Bob Hanson
Day 13 (Fri 3/8/2002):
Chapter 7, Sections 7.1-7.6
Entropy and the Second Law
Summary: Building on Chapter 4, this chapter emphasizes that change occurs in the
surroundings as well as in the system. Entropy is defined, and it is argued that the entropy
change of the surroundings is very important.
1. Entropy. The fact that energy is conserved argues that energy changes cannot
possibly “explain” chemical reactivity. Rather, the “most probable distribution”
has to include changes in the surroundings. The universe tends toward maximum
W for the simple reason that anything else is “less probable.”
2. The connection between ΔSsur and qsur. The result that ΔSsur = qsur /T is another
amazing property of Boltzmann distributions. Again, I wouldn’t do the proof in
class, but perhaps a little demo using WINTROPY.EXE would be in order. (They
do this in lab as part of Experiment 6.) For example, if you set up a system with
1000 particle and 1000 units of energy, you might get T = 104 Kelvin, W =
4.83E+592, ln W = 1364.646, and S = k ln W = 1.884E-20 J/K. (It may take some
time to find the most probable distribution. I think this is the “final” number, but
I’m not sure.) Then, adding 20 units of energy, so q = 20 and there are 1020 units
of energy, you might get T = 105 Kelvin and W = 3.75E+598, which gives ln W =
1378.267 and S = k ln W = 1.902E-20 J/K. Then ΔS is 1.88E-22 J/K. OK, now q =
20E-21 J (because each quantum unit is 1E-21 J) and the temperature is roughly
104.5 K. So q / T is approximately 1.91E-22 J/K. What do you know, it works!!!
3. Measuring entropy changes. The obvious problem is that heat usually results in a
change in temperature, so it’s not clear what “T” means in the equation ΔSsur =
qsur/T . The proper way to handle this is using calculus, but for our purposes we
can consider only changes where the temperature change is small.